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Pure chemistry 🧫⚗️💉 > Chapter 7 Part 1 : Mole concept > Flashcards

Chapter 7 Part 1 : Mole concept Flashcards

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1
Q

What makes up the mass of an atom?

A
  • The mass of an atom consists of mainly particles in the nucleus (protons and neutrons) since electrons have very small masses compared to both neutrons and protons.
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2
Q

Where can the mass of an atom of elements/nucleon numbers of an element be found?

A
  • The mass of an atom of most elements/nucleon numbers of an element, can be found in the periodic table.
  • This value corresponds to the relative atomic mass of the element.
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3
Q

How can we determine the relative atomic masses of elements in the periodic table?

A
  • The relative atomic masses of all elements in the periodic table are determined by comparing the masses of their atoms to 1/12 the mass of a carbon-12 atom.
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4
Q

What is the relative atomic mass (e.g. Ar)?

A
  • Relative atomic mass is defined as the average mass of one atom of an element compared to 1/12 the mass of a carbon-12 atom.
  • E.g. A sodium atom is 23 times heavier than 1/12 of a carbon-12 atom. Hence it has a relative atomic mass of 23.
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5
Q

Do atomic masses have any units?

A
  • No, this is because relative atomic masses is/are a ratio (a ratio does not have any units).
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6
Q

Are all of the relative atomic masses in the periodic table while numbers?

A
  • NOT all of the relative atomic masses in the periodic table are whole numbers because the presence of isotopes.
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7
Q

How is the relative atomic mass(es) of the elements determined?

A
  • The relative atomic masses of the elements in the periodic ta le determined using the average atomic masses of their isotopes.
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8
Q

What is the sum of the relative abundance of all isotopes of an element?

A
  • The sum of the relative abundance (ratio of the total number of species in an area to the total population of all species) of all isotopes of an element is 100%, and the relative atomic mass of the element is calculated using only the known abundances of the isotopes of the element.
  • E.g. Copper exists as 2 naturally occurring isotopes, copper-63 and copper-65, which have a natural abundance of 75% and 25% respectively.
  • Thus, the relative atomic mass of copper is (75/100 x 63) + (25/100 x 65) = 63.5

*** (Copper-63: natural abundance of 75%)
(Copper-65: natural abundance of 25%)

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9
Q

How is the mass of an element/compound calculated?

A
  • Some elements, and all compounds exist as more than one atom.
  • The mass of an element/a compound is calculated by adding up the relative atomic masses of all the atoms present in it.
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10
Q

What is the relative molecular mass?

A
  • The relative molecular mass (Mr) is defined as the mass of a substance compared to 1/12 the mass of a carbon-12 atom.
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11
Q

Does the relative molecular mass have any units?

A
  • No, just like relative atomic mass, relative molecular mass has no unit(s).
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12
Q

What is relative formula mass?

A
  • Relative formula mass is the mass of an ionic compound, and it is similar to relative molecular mass, and it can be represented by Mr.
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13
Q

What is the Mr of 2 substances?

A
  1. Fluorine
    Number of atoms in chemical formula:
    - 2F atoms
    Mr:
    - 2 x 19 (proton number) = 38
  2. Magnesium oxide (Mgo)
    Number of atoms in chemical formula:
    - 1 Mg atom
    - 1 O atom
    Mr:
    - 24 + 16 = 40
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14
Q

What is the percentage mass of an element in a compound?

A
  • The percentage mass of an element in a compound can be calculated using the expression:
  • (Number of atoms of elements x Ar of element/Mr of compound) x 100%
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15
Q

How can the percentage mass be calculated?

A
  • The percentage mass can be calculated by using the masses of the element and compound.
  • E.g. Mr of carbon dioxide:
    12 + (2 x 16) = 44
    Percentage mass of carbon in carbon dioxide:
    (12/44) x 100% = 27.3%
    Mass of carbon in 15g of carbon dioxide:
    27.3% x 15 = 4.10 g
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16
Q

What is the mole?

A
  • The mole is a form of measurement which is used to represent a fixed number of particles.
  • E.g. One mole of a substance represents 6.02 x 10(square 23) particles, and this number is the Avogadro constant.
  • Thus, one mole of different substance contains the same number of particles, and the particles can be atoms, molecules, ions, or electrons.
  • E.g. An ionic compound consists of ions, while a covalent substance consists of atoms/molecules.
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17
Q

What are the units of the mole?

A
  • The unit of the mole is Mol.
18
Q

How can the number of particles in a substance be calculated?

A
  • The number of particles in a substance can be calculated from the number of moles of the substance using the formula:
    Number of particles in a substance = number of moles of substance x 6.02 x 10(square 23)
19
Q

What are the 2 particles that are present in methane (CH4)?

A
  1. 1 mol of CH4
    Number of moles of particles:
    - 1 mol of CH4 molecules
    - 1 mol of C atoms
    - 4 mol of H atoms
    Number of particles:
    - 1 x 6.02 x 10(square 23) = 6.02 x 10(square 23) molecules.
    - 1 x 6.02 x 10(square 23) = 6.02 x 10(square 23) C atoms.
    - 4 x 6.02 x 10(square 23) = 2.41 x 10(square 24) H atoms.
  2. 2 mol of CH4:
    Number of moles of particles:
    - 2 mol of CH4 molecules
    - 2 mol of C4 atoms
    - 8 mol of H atoms
    Number of particles:
    - 2 x 6.02 x 10(square 23) = 1.20 x 10(square 24) CH4 molecules.
    - 2 x 6.02 x 10(square 23) = 1.20 x 10(square 24) C atoms.
    - 8 x 6.02 x 10(square 23) = 4.82 x 10(square 24) H atoms.
20
Q

How can the number of moles of a substance be calculated?

A
  • The number of moles of a substance can be calculated from the number of particles:
  • Number of moles of a substance = Number of particles of a substance/ 6.02 x 10(square 23)
21
Q

What is the molar mass of a substance?

A
  • The molar mass of a substance is equal to the relative atomic mass/relative molecular mass of 1 mole of the substance.
22
Q

What is the unit(s) of molar mass?

A
  • The unit of molar mass is g/mol.
23
Q

What can the molar mass of a substance be used to calculate?

A
  • The molar mass of a substance ca be used to calculate the number of moles of the substance/that the substance is using:
  • Number of moles of a substance = Mass of substance/molar mass of substance
24
Q

What are the number of moles of 2 substances with given masses?

A
  1. CO2 (carbon dioxide)
    Given mass/g:
    - 22g
    Molar mass/ g/mol:
    12 + (2 x 16)
    = 12 + 32
    = 44 g/mol
    Number of moles/mol:
    - 22/44 = 0.5
  2. Cu (NO3)2
    Given mass/g:
    - 20g
    Molar mass/ g/mol:
    64 + (2 x 14) + (2 x 3 x 16)
    = 64 + 28 + 96
    = 188 g/mol
    Number of moles/mol:
    - 20/188 = 0.106
25
Q

What can the (normal) molar mass be used to calculate?

A
  • The molar mass can be used to calculate the percentage mass of an element in a compound.
26
Q

What is the molar volume of a gas?

A
  • The molar volume of a gas is defined as the volume occupied by 1 mole of gas at room temperature, and pressure.
27
Q

does molar volume have a fixed value?

A
  • Yes, a molar volume has a fixed value of 24 dm3 and this means that 1 mole of any gas occupies a volume at 24 dm3 at room temperature, and pressure.
28
Q

How can we determine the number of moles of a gas at room temperature and pressure?

A
  • The number of moles of a gas at a room temperature and pressure can be determined through:
  • Number of moles = volume of gas in dm3/24 dm3
29
Q

How can the volume be converted from cm3 to dm3?

A
  • The volume can be converted from cm3 to dm3 based on he expression 1 dm3 = 1000cm3 = 100m3
30
Q

How can the mass of a gas relate to its volume?

A
  • The mass of a gas is related to its volume based on the expression:
  • Mass/molar mass = Volume/molar volume
31
Q

How is the concentration of a solution defined?

A
  • The concentration of a solution is defined as the amount of solute dissolved in a unit volume of the solvent.
32
Q

What are the units of concentration?

A
  • The units of concentration include g/dm3 and mol/dm3.
33
Q

What are the 2 examples of concentration?

A
  1. Copper (II) sulfate solution with a concentration of 20g/dm3 contains 20g of copper (II) sulfate dissolved in 1dm3 of water.
  2. Sodium chloride solution with a concentration of 0.500 mol/dm3 contains 0.500 mol of sodium chloride dissolved in 1 dm3 of water.
34
Q

How can the concentration of a solution be determined?

A

The concentration of a solution can be determined by:

  • Concentration in g/dm3 = Mass of solute in g/volume of solvent in dm3
  • Concentration in mol/dm3 = Number of moles of solute/volume of solvent in dm3
35
Q

How can the concentration in g/dm3 be converted into?

A

The concentration in g/dm3 can be converted into the concentration in mol/dm3 and vice versa:

  • Concentration in mol/dm3 = Concentration in g/dm3/molar mass
  • Concentration in g/dm3 = Concentration in mol/dm3 x molar mass

Example:
(E.g.) 32g of potassium bromide (Rbr) is dissolved in 150 cm3 of water.
- Molar mass of potassium bromide = 39+80=119g/mol
- Number of moles of potassium bromide
= 32/119
= 0.269 mol
- Concentration of potassium bromide in mol/dm3
= 0.269/ (150/1000)
= 1.79 mol/dm3
- Alternatively, concentration of potassium bromide in mol/dm3
= 213/119
= 1.79 mol/dm3

36
Q

What is an example of a molecular formula?

A
  • The molecular formula of carbon dioxide is CO2 (carbon dioxide).
  • This indicates that a carbon dioxide molecule is made up of 1 carbon atom and 2 oxygen atoms.
37
Q

What is a molecular formula?

A
  • The molecular formulae of a substance shows the exact number of atoms present in the substance.
38
Q

What is the empirical formula?

A
  • The empirical formula of a substance is the simplest form of the molecular formula.
  • It shows the ratio of elements that are present in the formula in its simplest form.
39
Q

What are the examples of molecular formulae(s) and their own empirical formulae?

A
  1. N2H4
    Mr of molecular formula:
    - 32
    Empirical formula:
    - NH2
    Mr of empirical formula:
    - 16
  2. C2H4O2
    Mr of molecular formula:
    - 60
    Empirical formula:
    - CH2O
    Mr of empirical formula:
    - 30
  3. C6H12O6
    Mr of molecular formula:
    - 180
    Empirical formula:
    - CH2O
    Mr of empirical formula:
    - 30
  4. P4O6
    Mr of molecular formula:
    - 220
    Empirical formula:
    - P2O3
    Mr of empirical formula:
    - 110
40
Q

How do we determine the multiple (molecular formulae is a multiple of the empirical formula)?

A
  • Since the molecular formula is a multiple of the epirical formula, we can determine the multiple by using:
  • (formula) n = Mr of molecular formula/Mr of empirical formula

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